Non-aqueous electrolytic secondary battery

ABSTRACT

A non-aqueous electrolyte secondary cell is provided with a positive electrode, a negative electrode, and a non-aqueous electrolyte solution, wherein said positive electrode comprises sulfur and said non-aqueous electrolyte solution comprises a room-temperature molten salt having a melting point of 60° C. or less.

TECHNICAL FIELD

The present invention relates to a non-aqueous electrolyte secondarycell provided with a positive electrode, a negative electrode, and anon-aqueous electrolyte solution, or more particularly, to a non-aqueouselectrolyte secondary cell enhanced in higher energy density bycombination of material employed as said positive electrode and thenon-aqueous electrolyte solution thereof.

BACKGROUND ART

As one type of secondary cells featuring high energy density, anon-aqueous electrolyte secondary cell has come into practical use,wherein a non-aqueous electrolyte solution is employed andcharge/discharge is performed by way of transferring lithium ionsbetween a positive electrode and a negative electrode.

Such a non-aqueous electrolyte secondary cell generally employs apositive electrode which is a lithium-transition metal compound oxidesuch as LiCoO₂ and the like, a negative electrode which is lithiummetal, a lithium alloy, or a carbon material capable of absorbing anddesorbing lithium, and a non-aqueous electrolyte solution wherein anelectrolyte of lithium salt, such as LiBF₄ or LiPF₆, is dissolved in anorganic solvent such as ethylene carbonate or diethyl carbonate.

Recently, further, such a non-aqueous electrolyte secondary cell hascome into practical use as an electric current source of portableequipment and the like, and accordingly, the non-aqueous electrolytesecondary cell having higher energy density has been desired.

Unfortunately, however, in such an ordinary non-aqueous electrolytesecondary cell, the lithium-transition metal compound oxide such asLiCoO₂ employed as the positive electrode thereof is large in weight andsmall in reactive electron number, therefore, capacity per unit weightis not sufficiently improved.

Further, sulfur is generally known as positive electrode material havinglarge theoretical capacity, however, where a simple substance of sulfuris employed as the positive electrode, very high temperature is requiredfor reversible reaction with lithium, therefore, the resultantnon-aqueous electrolyte secondary cell can not come into a generous use.

Therefore, recently, there has been proposed use of an organic disulfidecompound including DMcT (2,5-dimercapto-1,3,4-thiadiazole) as thepositive electrode material having large capacity and high energydensity. However, the organic disulfide compound used as the positiveelectrode material reversibly reacts with lithium only at a hightemperature of more than 60° C., therefore, the resultant non-aqueouselectrolyte secondary cell can not come into the generous use.

Further, more recently, Japanese Patent Application Nos.4-267073,8-115724, and so on have proposed the use of the positive electrodematerial which is a compound of the organic disulfide compound includingDMcT with conductive macromolecule including polyaniline for the purposeof charge/discharge reaction at normal temperature.

However, even in the use of the organic disulfide compound as thepositive electrode material, a part concerned with the charge/dischargereaction is only disulfide linkage portion, and other portions such ascarbon portion or hydrogen portion are not concerned with the reaction,therefore, the capacity per unit weight is not sufficiently improved.

DISCLOSURE OF THE INVENTION

The invention is directed to solution to aforementioned problems of anon-aqueous electrolyte secondary cell provided with a positiveelectrode, a negative electrode, and a non-aqueous electrolyte solution.

Specifically, the present invention has an object to provide anon-aqueous electrolyte secondary cell capable of charge/dischargereactions at normal temperature and having very high energy density,even in a case where the positive electrode comprises sulfur.

A non-aqueous electrolyte secondary cell according to a first aspect ofthe invention is provided with a positive electrode, a negativeelectrode, and a non-aqueous electrolyte solution, wherein said positiveelectrode comprises only a simple substance of sulfur as an activematerial and said non-aqueous electrolyte solution comprises aroom-temperature molten salt having a melting point of 60° C. or less.

A non-aqueous electrolyte secondary cell according to a second aspect ofthe invention is provided with a positive electrode, a negativeelectrode, and a non-aqueous electrolyte solution, wherein saidnon-aqueous electrolyte solution comprises a room-temperature moltensalt having a melting point of 60° C. or less and sulfur reductionproduct.

A non-aqueous electrolyte secondary cell according to a third aspect ofthe invention is provided with a positive electrode, a negativeelectrode, and a non-aqueous electrolyte solution, wherein said positiveelectrode comprises only a simple substance of sulfur as an activematerial and said non-aqueous electrolyte solution comprises aroom-temperature molten salt having a melting point of 60° C. or lessand at least one solvent selected from circular ether, chain ether, andcarbonate fluoride.

A non-aqueous electrolyte secondary cell according to a fifth aspect ofthe invention is provided with a positive electrode, a negativeelectrode using material capable of absorbing and desorbing lithium, anda non-aqueous electrolyte solution, wherein said positive electrodecomprises only a simple substance of sulfur as an active material andsaid non-aqueous electrolyte solution comprises a quaternary ammoniumsalt and a lithium salt.

As suggested by each of the above-mentioned non-aqueous electrolytesecondary cells, where the non-aqueous electrolyte secondary cell isprovided with the positive electrode comprising only the simplesubstance of sulfur as the active material, wherein the non-aqueouselectrolyte solution comprises any of the room-temperature molten salthaving the melting point of 60° C. or less, the room-temperature moltensalt having the melting point of 60° C. or less and lithium salt, theroom-temperature molten salt having the melting point of 60° C. or lessand at least one solvent selected from circular ether, chain ether, andcarbonate fluoride, or the quaternary ammonium salt and the lithiumsalt, sulfur in the positive electrode reversibly reacts with lithiumeven in normal temperature.

As suggested by the non-aqueous electrolyte secondary cell of the secondaspect of the invention, where the non-aqueous electrolyte solutioncomprises the room-temperature molten salt having the melting point of60° C. or less and sulfur reduction product, sulfur in the positiveelectrode reversibly reacts with lithium even at normal temperature, andthe non-aqueous electrolyte secondary cell can be charged/discharged atthe normal temperature, and in a case where the positive electrodecomprises sulfur, the non-aqueous electrolyte secondary cell also can becharged/discharged at the normal temperature.

The non-aqueous electrolyte solution can be used as a gel typeelectrolyte prepared by impregnating the above-mentioned non-aqueouselectrolyte solution with polymer electrolyte such as polyethylene oxideor poly acrylonitrile. The non-aqueous electrolyte solution can also beused as an inorganic solid electrolyte such as LiI or Li₃N.

In the non-aqueous electrolyte cells employing alkaline earth metals oralkaline metals except for lithium such as calcium, magnesium, sodium,or potassium, it is believed that sulfur in the positive electrodereversibly reacts with the alkaline metals or the alkaline earth metalsat the normal temperature for the effect of the above-mentionednon-aqueous electrolyte solution, thus, the non-aqueous electrolytesecondary cell can be charged/discharged at the normal temperature.

In the non-aqueous electrolyte secondary cell of the present invention,where the positive electrode comprises sulfur, the capacity per unitweight is further increased compared with organic disulfide.

In the positive electrode comprising sulfur, conductive agent ispreferably added to the positive electrode in order to improveconductivity, hereby to improve charge/discharge characteristics.Examples of the conductive agent include conductive carbon material. Inadding the conductive carbon material, an insufficient amount ofadditive results in insufficient improvement in the conductivity of thepositive electrode whereas an excessive amount of the additive resultsin insufficient ratio of sulfur in the positive electrode with theresult that large capacity is not attained. Therefore, the amount of thecarbon material based on the whole is generally set in a range of 5 to84 wt %, preferably in the range of 5 to 54 wt %, and more preferably inthe range of 5 to 20 wt %.

Where the room-temperature molten salt having the melting point of 60°C. or less is used as suggested by the non-aqueous electrolyte secondarycell of the invention, the aforesaid room-temperature molten salt is aliquid consisting only of ions, having no vapor pressure and havingnonflammable nature. Therefore, the room-temperature molten salt is notdecomposed nor caused to burn during the abnormal operation such asovercharge. That is, this salt ensures a safe use of the cell which isnot provided with the protection circuit. Where the lithium salt isadded to the room-temperature molten salt as described above, a meltingpoint of the resultant mixture is thought to be decreased from that ofeach of the salts so that the mixture is maintained in a liquid state.

What is required of the above room-temperature molten salt is to assumethe liquid state in a wide temperature range. In general, a usableroom-temperature molten salt may be in liquid state at temperatures inthe range of −20° C. to 60° C. and may preferably have a conductivity of10⁻⁴ S/cm or more.

Examples of a salt usable as such a room-temperature molten salt includequaternary ammonium salts and imidazolium salts.

Examples of usable quaternary ammonium salts as the room-temperaturemolten salt include at least one oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂,trimethyloctylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₈H₁₇)N⁻(SO₂CF₃)₂,trimethylallylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(Allyl)N⁻(SO₂CF₃)₂,trimethylhexylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₆H₁₃)N⁻(SO₂CF₃)₂,trimethylethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide(CH₃)₃N⁺(C₂H₅)(CF₃CO)N⁻(SO₂CF₃),trimethylallylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide(CH₃)₃N⁺(Allyl)(CF₃CO)N⁻(SO₂CF₃),trimethylpropylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide(CH₃)₃N⁺(C₃H₇)(CF₃CO)N⁻(SO₂CF₃),tetraethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide(C₂H₅)₄N⁺(CF₃CO)N⁻(SO₂CF₃),andtriethylmethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide(C₂H₅)₃N⁺(CH₃)(CF₃CO)N⁻(SO₂CF₃).

Examples of usable imidazolium salts as the room-temperature molten saltinclude at least one of1-ethyl-3-methylimidazolium.bis(pentafluoroethylsulfonyl)imide(C₂H₅)(C₃H₃N₂)⁺(CH₃)N⁻(SO₂C₂F₅)₂,1-ethyl-3-methylimidazolium.bis(trifluoromethylsulfonyl)imide(C₂H₅)(C₃H₃N₂)⁺(CH₃)N⁻(SO₂CF₃)₂,1-ethyl-3-methylimidazolium.tetrafluoroborate(C₂H₅)(C₃H₃N₂)⁺(CH₃)BF₄ ⁻,1-ethyl-3-methylimidazolium hexafluorophosphate(C₂H₅)(C₃H₃N₂)⁺(CH₃)PF₆⁻.

The non-aqueous electrolyte solution may comprise organic solvents suchas ethylene carbonate, diethyl carbonate, dimethyl carbonate, propylenecarbonate, cyclic ether, chain ether, and carbonate fluoride in additionto the room-temperature molten salt.

Examples of the usable cyclic ether include at least one of1,3-dioxolane, 4-methyl-1,3-dioxolane, tetrahydrofuran,2-methyltetrahydrofuran, propylene oxide, 1,2-butylene oxide,1,4-dioxane, 1,3,5-trioxane, furan, 2-methylfuran, 1,8-cineol, and crownether.

Examples of the usable chain ether include at least one of1,2-dimethoxyethane, diethyl ether, dipropyl ether, diisopropyl ether,dibuthyl ether, dihexyl ether, ethylvinyl ether, buthylvinyl ether,methylphenyl ether, ethylphenyl ether, buthylphenyl ether, pentylphenylether, methoxy toluene, benzylethyl ether, diphenyl ether, dibenzylether, o-dimethoxy benzene, 1,2-dimethoxyethane, 1,2-dibutoxyethane,diethylene glycol dimethyl ether, diethylene glycol ethyl ether,diethylene glycol dibuthyl ether, 1,1-dimethoxy methane, 1,1-diethoxyethane, triethylene glycol dimethyl ether, and tetraethylene glycoldimethyl ether.

Examples of the usable carbonate fluoride include at least one oftrifluoropropylene carbonate and fluoroethylene carbonate.

Examples of the quaternary ammonium salts usable for the non-aqueouselectrolyte solution of the non-aqueous electrolyte secondary cellaccording to the fifth aspect of the invention includetetramethylammonium.tetrafluoroborate(CH₃)₄N⁺BF₄,tetramethylammonium.hexafluorophosphate(CH₃)₄N⁺PF₆,tetraethylammonium.tetrafluoroborate(C₂H₅)₄N⁺BF₄, andtetraethylammonium.hexafluorophosphate (C₂H₅)₄N⁺PF₆ in addition to thequaternary ammonium salts as the room-temperature molten salt.

Electrolytes commonly used in the conventional non-aqueous electrolytesecondary cells may be used as the lithium salt to be mixed with thenon-aqueous electrolyte solution of the non-aqueous electrolytesecondary cell of the present invention. Examples of the usable lithiumsalt include at least of LiBF₄, LiPF₆, LiCF₃SO₃, LiC₄F₉SO₃,LiN(CF₃SO₂)₂, LiN(C₂F₅SO₂)₂, LiN(CF₃SO₂)(COCF₃) and LiAsF₆.

In the non-aqueous electrolyte secondary cell of the present invention,materials commonly used in the conventional non-aqueous electrolytesecondary cells may be used as the material capable of absorbing anddesorbing lithium for the negative electrode thereof. Examples of suchmaterials include carbon materials such as lithium metal, lithium alloy,and graphite, however, for the purpose of obtaining the non-aqueouselectrolyte secondary cell of high energy density, it is particularlydesirable to employ silicon having large capacity as suggested inJapanese Patent Application Nos.2000-321200 and 2000-321201 filed by thepresent applicant.

In the non-aqueous electrolyte secondary cell of the present invention,one of the positive electrode and the negative electrode is providedwith lithium which is concerned with the charge/discharge reactions.

BRIEF DESCRIPTION OF THE DRAWINGS

FIG. 1 is a diagram schematically illustrating a test cell fabricated inExamples 1 to 20 and comparative examples 1 to 5 of the invention;

FIG. 2 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 1;

FIG. 3 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of comparative example 1;

FIG. 4 is a graph representing initial charge/discharge characteristicsof the test cell of Example 1;

FIG. 5 is a graph representing discharge capacity and charge-dischargeefficiency in each cycle of repeated charge/discharges of the test cellof Example 1;

FIG. 6 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of reference example 1;

FIG. 7 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 3;

FIG. 8 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of comparative example 2;

FIG. 9 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 4;

FIG. 10 is a graph representing initial charge/discharge characteristicsof the test cell of Example 4;

FIG. 11 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of a positive electrode of thetest cell of Example 5;

FIG. 12 is a graph representing cyclic voltammetry measured by scanningpotential of the positive electrode of the test cell of Example 6;

FIG. 13 is a graph representing initial charge/discharge characteristicsof the test cell of Example 6;

FIG. 14 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 7;

FIG. 15 is a graph representing initial charge/discharge characteristicsof the test cell of Example 7;

FIG. 16 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 8;

FIG. 17 is a graph representing initial charge/discharge characteristicsof the test cell of Example 8;

FIG. 18 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of reference example 2;

FIG. 19 is a graph representing initial charge/discharge characteristicsof the test cell of reference example 2;

FIG. 20 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 10;

FIG. 21 is a graph representing initial charge/discharge characteristicsof the test cell of Example 10;

FIG. 22 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 11;

FIG. 23 is a graph representing initial charge/discharge characteristicsof the test cell of Example 11;

FIG. 24 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of comparative example 3;

FIG. 25 is a graph representing initial charge/discharge characteristicsof the test cell of comparative example 3;

FIG. 26 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 12;

FIG. 27 is a graph representing initial charge/discharge characteristicsof the test cell of Example 12;

FIG. 28 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 13;

FIG. 29 is a graph representing initial charge/discharge characteristicsof the test cell of Example 13;

FIG. 30 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of comparative example 4;

FIG. 31 is a graph representing initial charge/discharge characteristicsof the test cell of comparative example 4;

FIG. 32 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 14;

FIG. 33 is a graph representing initial charge/discharge characteristicsof the test cell of Example 14;

FIG. 34 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 15;

FIG. 35 is a graph representing initial charge/discharge characteristicsof the test cell of Example 15;

FIG. 36 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of comparative example 5;

FIG. 37 is a graph representing initial charge/discharge characteristicsof the test cell of comparative example 5;

FIG. 38 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 16;

FIG. 39 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 17;

FIG. 40 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 18;

FIG. 41 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 19;

FIG. 42 is a graph representing initial charge/discharge characteristicsof the test cell of Example 19; and

FIG. 43 is a graph representing cyclic voltammetry of a positiveelectrode measured by scanning potential of the positive electrode ofthe test cell of Example 20.

BEST MODES FOR CARRYING OUT THE INVENTION

Examples will make it apparent that non-aqueous electrolyte secondarycells of the present invention are properly charged/discharged at roomtemperature, and have very high energy density, even in a case where apositive electrode comprises sulfur. It is to be distinctly appreciatedthat the non-aqueous electrolyte secondary cells of the presentinvention should not be limited to the following examples butappropriate changes and modifications may be made in carrying out theinvention without departing from the spirit and scope of the invention.

EXAMPLE 1

Example 1 used a non-aqueous electrolyte solution prepared by dissolvingof LiN(CF₃SO₂)₂ as a lithium salt intrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂as a room-temperature molten salt in a concentration of 0.3 mol/l.

A positive electrode was prepared as follows. 20 parts by weight ofsulfur, 70 parts by weight of acetylene black as a conductive agent, and10 parts by weight of polytetrafluoroethylene as a binder were kneaded,then resultant mixture was stirred in a mortar for 30 minutes, and wassubject to a pressure of 150 kg/cm² for 5 seconds in a molding die to beformed into a disk with a diameter of 10.3 mm. The disk was covered withan aluminum net.

A test cell of Example 1 was fabricated as follows. As shown in FIG. 1,the aforesaid non-aqueous electrolyte solution 14 was poured into a testcell vessel 10. On the other hand, the aforesaid positive electrode 11was used as a working electrode whereas lithium metal pieces were usedas a negative electrode 12 as a counter electrode and a referenceelectrode 13.

COMPARATIVE EXAMPLE 1

Comparative example 1 used a non-aqueous electrolyte solution preparedby mixing ethylene carbonate and dimethyl carbonate in a volume ratio of1:1 to give a mixture solvent, in which LiPF₆ as a lithium salt wasdissolved in a concentration of 1 mol/l. Except for the above, the sameprocedure as that in Example 1 was taken to fabricate a test cell ofcomparative example 1.

Cyclic volutammetry of the test cell of Example 1 was determined asfollows. Potential scanning range of the positive electrode 11 versusthe reference electrode 13 was set to be in a range of 1 to 5 V(vs.Li/Li⁺), while potential scanning rate was set to be 0.5 mV/s. Then2 cycles of an operation comprising steps of scanning an initialpotential of the positive electrode 11, 2.9 V (vs.Li/Li⁺), versus thereference electrode 13 in a reduction direction, and scanning saidinitial potential in an oxidation direction were carried out so as todetermine the cyclic volutammetry at each cycle. The results are shownin FIG. 2.

In addition, the cyclic volutammetry of the test cell of comparativeexample 1 was determined as follows. Potential scanning range of thepositive electrode 11 versus the reference electrode 13 was set to be ina range of 1 to 4.2 V (vs.Li/Li⁺), while potential scanning rate was setto be 0.5 mV/s. Then 2 cycles of the operation comprising steps ofscanning the initial potential of the positive electrode 11, 3.0V(vs.Li/Li⁺), versus the reference electrode 13 in the reductiondirection, and scanning said initial potential in the oxidationdirection were carried out so as to determine the cyclic volutammetry ateach cycle. The results are shown in FIG. 3.

As a result, in the test cell of Example 1, in the scanning in thereduction direction, the reduction current began to extremely flowaround less than 2.3 V (vs.Li/Li⁺), and sulfur was expectedly reduced.In the scanning in the oxidation direction, oxidation peak existed in arange of 2.6 to 3.9 V (vs.Li/Li⁺), and the reduced sulfur was expectedlyoxidized in said potential range. The same results were attained also inthe second cycle, and it is believed that sulfur reversibly reacted.

On the other hand, in the test cell of comparative example 1, in thescanning in the reduction direction, the reduction current began to flowaround less than 2.4 V (vs.Li/Li⁺), and sulfur was expectedly reduced.In the scanning in the oxidation direction, however, the oxidation peakdid not exist, and the reduced sulfur expectedly was not oxidized. Inthe second cycle, in the scanning in the reduction direction, thereduction current flew around less than 2.4 V (vs.Li/Li⁺), expectedlybecause the sulfur which was not reduced and remained in the priorreaction was reduced this time.

The test cell of Example 1 was subject to a discharge at dischargingcurrent of 0.13 mA/cm² to a discharge cut-off potential of 1 V(vs.Li/Li⁺) and a charge at charging current of 0.13 mA/cm² to a chargecut-off potential of 2.7V (vs.Li/Li⁺), so as to inspect an initialcharge/discharge characteristics thereof. The results are shown in FIG.4. A discharge curve showing a relation between the potential in thedischarge and capacity density per gram of sulfur is represented by asolid line, whereas a charge curve showing a relation between thepotential in the charge and the capacity density per gram of sulfur isrepresented by a broken line.

As a result, in the test cell of Example 1, an initial dischargecapacity density per gram of sulfur was about 654 mAh/g, which was lessthan 1675 mAh/g according to theoretical capacity density but wasextremely higher than that of LiCoO₂ which has been generally used inthe positive electrode. In addition, an initial charge capacity densityper gram of sulfur was so high as about 623 mAh/g, which showed thatsulfur reversibly reacted.

Further, the test cell of Example 1 was subject to repeatedcharge/discharge processes comprising the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1 V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 2.7V (vs.Li/Li⁺), so as to determine acharge capacity Qa (mAh/g) and a discharge capacity Qb (mAh/g) of thecell. Charge/discharge efficiency (%) in each cycle was determined basedon the following equation. In FIG. 5, the discharge capacity (mAh/g) ineach cycle is represented by a combination of a hollow circle and asolid line, whereas the charge/discharge efficiency (%) in each cycle isrepresented by a combination of a triangle and a broken line.Charge/discharge efficiency (%)=(Qb/Qa)×100

The results show that the test cell of Example 1 stably had thedischarge capacity of about 490 mAh/g in the third and the succeedingcycles, and the charge/discharge efficiency was stably 100%.

In the test cell of Example 1, an average discharge voltage was about 2V, and an energy density per gram of sulfur was about 980 mWh/g whichwas higher than that of LiCoO₂ (about 540 mWh/g) which has beengenerally used in the positive electrode.

REFERENCE EXAMPLE 1

Reference example 1 used the non-aqueous electrolyte solution preparedby dissolving LiPF₆ in γ-butyrolactone in a concentration of 1 mol/l.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of reference example 1.

In addition, the cyclic volutammetry of the test cell of referenceexample 1 was determined as follows. The potential scanning range of thepositive electrode 11 versus the reference electrode 13 was set to be ina range of 1 to 4.7 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then 3 cycles of the operation comprising steps ofscanning the initial potential of the positive electrode 11, 3.34(vs.Li/Li⁺), versus the reference electrode 13 in the reductiondirection, and scanning said initial potential in the oxidationdirection were carried out so as to determine the cyclic volutammetry ateach cycle. The results are shown in FIG. 6.

As a result, in the test cell of reference example 1, in the scanning inthe reduction direction, the reduction current began to flow around lessthan 2.3 V (vs.Li/Li⁺), and sulfur was expectedly reduced. In thescanning in the oxidation direction, the oxidation peak existed in therange of 2.5 to 3.6V (vs.Li/Li⁺), and the reduced sulfur was expectedlyoxidized in said potential range. The same results were attained also inthe second cycle, and it is believed that sulfur reversibly reacted. Thepresent reference example used γ-butyrolactone as the solvent of thenon-aqueous electrolyte solution, however, roughly the same results maybe attained by use of circular ester including γ-valerolactone otherthan γ-butyrolactone.

EXAMPLE 3

Example 3 used the non-aqueous electrolyte solution prepared bydissolving LiPF₆ as the lithium salt in a concentration of 1 mol/l in amixture solvent containing tetrafluoro propylene carbonate which iscarbonate fluoride andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂as a quaternary ammonium salt in a volume ratio of 1:1. Except for theabove, the same procedure as that in Example 1 was taken to fabricate atest cell of Example 3.

COMPARATIVE EXAMPLE 2

Comparative example 2 used the non-aqueous electrolyte solution preparedby dissolving LiPF₆ as the lithium salt in tetrafluoro propylenecarbonate in the concentration of 1 mol/l. Except for the above, thesame procedure as that in Example 1 was taken to fabricate a test cellof comparative example 2.

In addition, the cyclic volutammetry of the test cells of Example 3 andcomparative example 2 was determined as follows. The potential scanningrange of the positive electrode 11 versus the reference electrode 13 wasset to be in the range of 1 to 4.7 V (vs.Li/Li⁺), while the potentialscanning rate was set to be 1 mV/s. Then the operation comprising stepsof scanning the initial potential of the positive electrode 11, 3.34 V(vs.Li/Li⁺), versus the reference electrode 13 in the reductiondirection, and scanning said potential in the oxidation direction wascarried out for 4 cycles to the test cell of Example 3 and for 3 cyclesto the test cell of comparative example 2, so as to determine the cyclicvolutammetry at each cycle. The results of the test cell of Example 3are shown in FIG. 7, and the results of the test cell of comparativeexample 2 are shown in FIG. 8.

As a result, in the test cell of Example 3, in the scanning in thereduction direction, the reduction current began to flow around lessthan 2.3 V (vs.Li/Li⁺), and sulfur is expectedly reduced. In thescanning in the oxidation direction, oxidation peak existed in the rangeof 2.0 to 3.0 V (vs.Li/Li⁺), and the reduced sulfur was expectedlyoxidized in said potential range. The same results were attained also inthe second cycle, and it is believed that sulfur reversibly reacted.

On the other hand, in the test cell of comparative example 2, in thescanning in the reduction direction, the reduction current began to flowaround less than 2.2 V (vs.Li/Li⁺), and sulfur was expectedly reduced.In the scanning in the oxidation direction, oxidation peak existedaround 4V (vs.Li/Li⁺), and energy efficiency was degraded. From thesecond cycle and onwards, the oxidation peak and the reduction currentwere steeply decreased, and the reversibility was reduced.

In the test cell of Example 3, discharge potential of sulfur was about2.0 V (vs. Li/Li⁺), and the energy density converted from thetheoretical capacity density of sulfur, 1675 mAh/g, was 3350 mWh/g,which is extremely higher than that of LiCoO₂ (about 540 mWh/g) whichhas been generally used in the positive electrode.

EXAMPLE 4

Example 4 used the non-aqueous electrolyte solution prepared bydissolving LiN(CF₃SO₂)₂ as the lithium salt intrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂as the room-temperature molten salt in the concentration of 0.3 mol/l.Then, in the non-aqueous electrolyte solution, sulfur and lithium weremade to contact with each other to synthesize reduction product ofsulfur. Further, the sulfur and lithium which were unreacted werewithdrawn so as to prepare the non-aqueous electrolyte solutioncontaining the reduction product of sulfur.

Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 4.

In addition, the cyclic volutammetry of the test cell of Example 4 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1 to 4.7 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.7 V (vs.Li/Li⁺),versus the reference electrode 13 in the oxidation direction, andscanning said potential in the reduction direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 9.

As a result, in the test cell of Example 4, in the first scanning in theoxidation direction, the peak corresponding to the oxidation of sulfurthus reduced did not appear, however, in the scanning in the reductiondirection, the reduction current began to flow around less than 2.4 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, oxidation peak existed in the range of 2.2 to 3.9 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized in saidpotential range. The same results were attained also from the secondcycle and onwards, and it is believed that sulfur reversibly reacted.

The test cell of Example 4 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1 V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 3.5V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 10. The discharge curve showing the relation between thepotential in the discharge and the capacity density per gram of sulfuris represented by the solid line, whereas the charge curve showing therelation between the potential in the charge and the capacity densityper gram of sulfur is represented by the broken line.

As a result, in the test cell of Example 4, the initial dischargecapacity density per gram of sulfur was about 749 mAh/g, which was lessthan 1675 mAh/g according to theoretical capacity density but wasextremely higher than that of LiCoO₂ which has been generally used inthe positive electrode. In addition, the discharge capacity density ofthe test cell of Example 4 containing the reduction product of sulfur inthe non-aqueous electrolyte solution thereof was higher than that of thetest cell of Example 1.

EXAMPLE 5

Example 5 used the non-aqueous electrolyte solution prepared in the samemanner as the Example 4.

The positive electrode was prepared as follows. 80 parts by weight ofacetylene black as the conductive agent and 20 parts by weight ofpolytetrafluoro ethylene as the binder were kneaded, then the resultantmixture was stirred in the mortar for 30 minutes, and was subject to thepressure of 150 kg/cm² for 5 seconds in the molding die to be formedinto the disk with the diameter of 10.3 mm. The disk was covered withthe aluminum net.

Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 5.

In addition, the cyclic volutammetry of the test cell of Example 5 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1 to 4.7 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.3 V (vs.Li/Li⁺),versus the reference electrode 13 in the oxidation direction, andscanning said potential in the reduction direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 11.

As a result, in the test cell of Example 5, in the first scanning in theoxidation direction, the peak corresponding to the oxidation of sulfurthus reduced hardly appeared, however, in the scanning in the reductiondirection, the reduction current began to flow around less than 2.3 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak existed in the range of 2.3 to3.4V (vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized in saidpotential range. The same results were attained also from the secondcycle and onwards, and it is believed that sulfur reversibly reacted.

EXAMPLE 6

Example 6 usedtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂which is the room-temperature molten salt as the non-aqueous electrolytesolution without adding the lithium salt. Except for the above, the sameprocedure as that in Example 1 was taken to fabricate a test cell ofExample 6.

In addition, the cyclic volutammetry of the test cell of Example 6 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1 to 4.7 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.3 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 12.

As a result, in the test cell of Example 6, in the scanning in thereduction direction, the reduction current began to flow around lessthan 2 V (vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanningin the oxidation direction, the oxidation peak existed around 4 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential. The same results were attained also from the second cycle andonwards, and it is believed that sulfur reversibly reacted.

The test cell of Example 6 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1 V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 4.5V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 13. The discharge curve showing the relation between thepotential in the discharge and the capacity density per gram of sulfurwas represented by the solid line, whereas the charge curve showing therelation between the potential in the charge and the capacity densityper gram of sulfur is represented by the broken line.

As a result, in the test cell of Example 6, the initial dischargecapacity density per gram of sulfur was about 366 mAh/g, which was lessthan 1675 mAh/g according to theoretical capacity density but wasextremely higher than that of LiCoO₂ which has been generally used inthe positive electrode.

The results of Example 6 show that sulfur is charged/discharged evenwithout presence of the lithium salt in the non-aqueous electrolytesolution. Therefore, not only where the negative electrode comprisesmaterial capable of absorbing and desorbing lithium ions but also wherethe negative electrode comprises material capable of absorbing anddesorbing cation, sulfur is charged/discharged. Examples of said cationinclude alkaline earth metal ions such as calcium ion or magnesium ion,and alkaline metal ions such as sodium ion or potassium ion. Thenon-aqueous electrolyte solution may comprise alkaline earth salts suchas calcium salt or magnesium salt, and alkaline metal salts such assodium salt or potassium salt.

EXAMPLE 7

Example 7 used the non-aqueous electrolyte solution prepared bydissolving LiN(CF₃SO₂)₂ as the lithium salt intriethylmethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide(C₂H₅)₃N⁺(CH₃)(CF₃CO)N⁻(SO₂CF₃)as the room-temperature molten salt in a concentration of 0.5 mol/l.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 7.

In addition, the cyclic volutammetry of the test cell of Example 7 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1 to 4.7 V (vs.Li/Li⁺), while potential scanning rate was setto be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 3.0 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 14.

As a result, in the test cell of Example 7, in the scanning in thereduction direction, the reduction current began to flow around lessthan 2.3 V (vs.Li/Li⁺), and sulfur was expectedly reduced. In thescanning in the oxidation direction, the oxidation peak existed around3.8 V (vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized aroundsaid potential. The same results are attained also from the second cycleand onwards, and it is believed that sulfur reversibly reacted.

The test cell of Example 7 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1 V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 3.5V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 15. The discharge curve showing the relation between thepotential in the discharge and the capacity density per gram of sulfuris represented by the solid line, whereas the charge curve showing therelation between the potential in the charge and the capacity densityper gram of sulfur is represented by the broken line.

As a result, in the test cell of Example 7, the initial dischargecapacity density per gram of sulfur was 1138 mAh/g, which was extremelyhigher than that of LiCoO₂ which has been generally used in the positiveelectrode.

EXAMPLE 8

Example 8 used the non-aqueous electrolyte solution prepared bydissolving LiN(CF₃SO₂)₂ as the lithium salt intrimethylhexylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₆H₁₃)N⁻(SO₂CF₃)₂as the room-temperature molten salt in a concentration of 0.5 mol/l.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 8.

In addition, the cyclic volutammetry of the test cell of Example 8 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1 to 4.7 V (vs.Li/Li⁺), while potential scanning rate was setto be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.8 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 16.

As a result, in the test cell of Example 8, in the scanning in thereduction direction, the reduction current began to flow around lessthan 2.3 V (vs.Li/Li⁺), and sulfur was expectedly reduced. In thescanning in the oxidation direction, the oxidation peak existed around2.6 V (vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized aroundsaid potential. The same results were attained also from the secondcycle and onwards, and it is believed that sulfur reversibly reacted.

The test cell of Example 8 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1.0V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 3.5V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 17. The discharge curve showing the relation between thepotential in the discharge and the capacity density per gram of sulfuris represented by the solid line, whereas the charge curve showing therelation between the potential in the charge and the capacity densityper gram of sulfur is represented by the broken line.

As a result, in the test cell of Example 8, the initial dischargecapacity density per gram of sulfur was 588 mAh/g, which was extremelyhigher than that of LiCoO₂ which has been generally used in the positiveelectrode.

REFERENCE EXAMPLE 2

Reference example 2 used the positive electrode prepared as follows. 90parts by weight of copper sulfide CuS, 5 parts by weight of acetyleneblack as the conductive agent, and 5 parts by weight of polytetrafluoroethylene as the binder were kneaded, then the resultant mixture wasstirred in the mortar for 30 minutes, and was subject to the pressure of150 kg/cm² for 5 seconds in the molding die to be formed into the diskwith the diameter of 10.3 mm. The disk was covered with the aluminumnet. Except for the above, the same procedure as that in Example 1 wastaken to fabricate a test cell of reference example 2.

In addition, the cyclic volutammetry of the test cell of referenceexample 2 was determined as follows. The potential scanning range of thepositive electrode 11 versus the reference electrode 13 was set to be inthe range of 1 to 3.7 V (vs.Li/Li⁺), while potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.3 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 18.

As a result, in the test cell of reference example 2, in the scanning inthe reduction direction, the reduction current began to flow around lessthan 2.3 V (vs.Li/Li⁺), and copper sulfide was expectedly reduced. Inthe scanning in the oxidation direction, the oxidation peak existedaround 2.8 V (vs.Li/Li⁺), and the reduced copper sulfide was expectedlyoxidized around said potential. The same results were attained also fromthe second cycle and onwards, and it is believed that copper sulfidereversibly reacted.

The test cell of reference example 2 was subject to the discharge atdischarging current of 0.13 mA/cm² to the discharge cut-off potential of1.0 V (vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² tothe charge cut-off potential of 2.7 V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 19. The discharge curve showing the relation between thepotential in the discharge and the capacity density per gram of coppersulfide is represented by the solid line, whereas the charge curveshowing the relation between the potential in the charge and thecapacity density per gram of copper sulfide is represented by the brokenline.

As a result, in the test cell of reference example 2, the initialdischarge capacity density per gram of copper sulfide was 129 mAh/g.

EXAMPLE 10

Example 10 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in the mixturecontaining 50% by volume of 1,3-dioxolane and 50% by volume oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 10.

In addition, the cyclic volutammetry of the test cell of Example 10 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.4 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 20.

As a result, in the test cell of Example 10, in the scanning in thereduction direction, the reduction current began to flow around lessthan 2.3 V (vs.Li/Li⁺), and sulfur was expectedly reduced. In thescanning in the oxidation direction, the oxidation peak existed around2.6V (vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized aroundsaid potential. The same results are attained also from the second cycleand onwards, and it is believed that sulfur reversibly reacted.

The test cell of Example 10 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1.0 V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 3.0V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 21. The discharge curve showing the relation between thepotential in the discharge and the capacity density per gram of sulfuris represented by the solid line, whereas the charge curve showing therelation between the potential in the charge and the capacity densityper gram of sulfur is represented by the broken line.

As a result, in the test cell of Example 10, the initial dischargecapacity density per gram of sulfur was 2230 mAh/g, which was extremelyhigher than that of LiCoO₂ which has been generally used in the positiveelectrode. In addition, where both 1,3-dioxolane andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂were mixed, the discharge capacity density was higher around more than2.0 V (vs.Li/Li⁺) compared with the case in which the non-aqueouselectrolyte solution comprised only 1,3-dioxolane as the solvent thereofas shown in the following comparative example 3, and the dischargecapacity density was also higher compared with the case in which thenon-aqueous electrolyte solution comprised onlytrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂as the solvent thereof as shown in Example 1.

EXAMPLE 11

Example 11 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in the mixturecontaining 25% by volume of 1,3-dioxolane and 75 by volume oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 11.

In addition, the cyclic volutammetry of the test cell of Example 11 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.3 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.4 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 22.

As a result, in the test cell of Example 11, in the scanning in thereduction direction, the reduction peak appeared around 1.9 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 2.4 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential. From the second cycle and onwards, in the scanning in thereduction direction, the reduction peak existed around 1.5 V(vs.Li/Li⁺), and in the scanning in the oxidation direction, theoxidation peak existed around 2.4 V (vs.Li/Li⁺), thus, it is believedthat sulfur reversibly reacted.

The test cell of Example 11 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1.0 V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 3.0V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 23. The discharge curve showing the relation between thepotential in the discharge and capacity density per gram of sulfur isrepresented by the solid line, whereas the charge curve showing therelation between the potential in the charge and capacity density pergram of sulfur is represented by the broken line.

As a result, in the test cell of Example 11, the initial dischargecapacity density per gram of sulfur was 2291 mAh/g, which was extremelyhigher than that of LiCoO₂ which has been generally used in the positiveelectrode. In addition, where both 1,3-dioxolane andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂were mixed, the discharge capacity density was higher around more than2.0 V (vs.Li/Li⁺) compared with the case in which the non-aqueouselectrolyte solution comprised only 1,3-dioxolane as the solvent thereofas shown in the following comparative example 3, and the dischargecapacity density was also higher compared with the case in which thenon-aqueous electrolyte solution comprised onlytrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂as the solvent thereof as shown in Example 1.

COMPARATIVE EXAMPLE 3

Comparative example 3 used the non-aqueous electrolyte solution preparedby dissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in1,3-dioxolane. Except for the above, the same procedure as that inExample 1 was taken to fabricate a test cell of comparative example 3.

In addition, the cyclic volutammetry of the test cell of comparativeexample 3 was determined as follows. The potential scanning range of thepositive electrode 11 versus the reference electrode 13 was set to be inthe range of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning ratewas set to be 1.0 mV/s. Then the operation comprising steps of scanningthe initial potential of the positive electrode 11, 2.2 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 24.

As a result, in the test cell of comparative example 3, in the scanningin the reduction direction, the reduction peak appeared around 1.8 V(vs. Li/Li⁺) and the high reduction current began to flow around lessthan 1.2 V (vs.Li/Li⁺), and sulfur was expectedly reduced. In thescanning in the oxidation direction, the oxidation peak existed around2.6 V (vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized aroundsaid potential.

The test cell of comparative example 3 was subject to the discharge atdischarging current of 0.13 mA/cm² to the discharge cut-off potential of1.0 V (vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² tothe charge cut-off potential of 3.0 V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 25. The discharge curve showing the relation between thepotential in the discharge and capacity density per gram of sulfur isrepresented by the solid line, whereas the charge curve showing therelation between the potential in the charge and capacity density pergram of sulfur is represented by the broken line.

As a result, in the test cell of comparative example 3, the initialdischarge capacity density per gram of sulfur was 1677 mAh/g, which wasextremely higher than that of LiCoO₂ which has been generally used inthe positive electrode, however, the discharge potential thereof was solow as about 1.2 V (vs.Li/Li⁺).

Example 10 and Example 11 show that wheretrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂and 1,3-dioxolane were mixed together, viscosity of the non-aqueouselectrolyte solution was less, which was more preferable compared withthe case in which onlytrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂was used. Further, the results of Example 1, Example 10, Example 11, andcomparative example 3 show that where the positive electrode comprisedsulfur, the use of the mixture oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂and 1,3-dioxolane was more preferable than the sole use oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂or 1,3-dioxolane, according to comparison of each discharge capacitydensity around more than 2 V (vs.Li/Li⁺) in the discharge. The amount of1,3-dioxolane therein is generally set in a range of 0.1 to 99.9% byvolume, preferably in the range of 0.1 to 50% by volume, and morepreferably in the range of 0.1 to 25% by volume.

EXAMPLE 12

Example 12 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in the mixturecontaining 50% by volume of tetrahydrofuran and 50% by volume oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 12.

In addition, the cyclic volutammetry of the test cell of Example 12 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.5 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 26.

As a result, in the test cell of Example 12, in the scanning in thereduction direction, the reduction peak appeared both around 2.0 V(vs.Li/Li⁺) and 1.5 V (vs.Li/Li⁺), and sulfur was expectedly reduced. Inthe scanning in the oxidation direction, the oxidation current flewaround more than 2.2 V (vs.Li/Li⁺), and the reduced sulfur wasexpectedly oxidized more than said potential.

The test cell of Example 12 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1.0 V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 3.0V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 27. The discharge curve showing the relation between thepotential in the discharge and capacity density per gram of sulfur isrepresented by the solid line, whereas the charge curve showing therelation between the potential in the charge and capacity density pergram of sulfur is represented by the broken line.

As a result, in the test cell of Example 12, the initial dischargecapacity density per gram of sulfur was 1479 mAh/g, which was extremelyhigher than that of LiCoO₂ which has been generally used in the positiveelectrode. In addition, where both tetrahydrofuran andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂were mixed, the discharge capacity density was higher around more than2.0 V (vs.Li/Li⁺) compared with the case in which the non-aqueouselectrolyte solution comprised only tetrahydrofuran as the solventthereof as shown in the following comparative example 4, and thedischarge capacity density was also higher compared with the case inwhich the non-aqueous electrolyte solution comprised onlytrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂as the solvent thereof as shown in Example 1.

EXAMPLE 13

Example 13 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂) as the lithium salt in the mixturecontaining 25% by volume of tetrahydrofuran and 75 by volume oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 13.

In addition, the cyclic volutammetry of the test cell of Example 13 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.6 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 28.

As a result, in the test cell of Example 13, in the scanning in thereduction direction, the reduction current flew around less than 2.4 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 2.5 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential.

The test cell of Example 13 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1.0 V (vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to the chargecut-off potential of 3.0 V (vs.Li/Li⁺), so as to inspect the initialcharge/discharge characteristics thereof. The results are shown in FIG.29. The discharge curve showing the relation between the potential inthe discharge and capacity density per gram of sulfur is represented bythe solid line, whereas the charge curve showing the relation betweenthe potential in the charge and capacity density per gram of sulfur isrepresented by the broken line.

As a result, in the test cell of Example 13, the initial dischargecapacity density per gram of sulfur was 1547 mAh/g, which was extremelyhigher than that of LiCoO₂ which has been generally used in the positiveelectrode. In addition, where both tetrahydrofuran andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂were mixed, the discharge capacity density was higher around more than2.0 V (vs.Li/Li⁺) compared with the case in which the non-aqueouselectrolyte solution comprised only tetrahydrofuran as the solventthereof as shown in the following comparative example 4, and thedischarge capacity density was also higher compared with the case inwhich the non-aqueous electrolyte solution comprised onlytrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂as the solvent thereof as shown in Example 1.

COMPARATIVE EXAMPLE 4

Comparative example 4 used the non-aqueous electrolyte solution preparedby dissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt intetrahydrofuran. Except for the above, the same procedure as that inExample 1 was taken to fabricate a test cell of comparative example 4.

In addition, the cyclic volutammetry of the test cell of comparativeexample 4 was determined as follows. The potential scanning range of thepositive electrode 11 versus the reference electrode 13 was set to be inthe range of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning ratewas set to be 1.0 mV/s. Then the operation comprising steps of scanningthe initial potential of the positive electrode 11, 2.3 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 30.

As a result, in the test cell of comparative example 4, in the scanningin the reduction direction, the reduction peak appeared around 1.6 V(vs.Li/Li⁺) and the high reduction current flew around less than 1.2 V(vs.Li/Li⁺), thus sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 2.5 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential.

The test cell of comparative example 4 was subject to the discharge atdischarging current of 0.13 mA/cm² to the discharge cut-off potential of1.0 V (vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² tothe charge cut-off potential of 3.3 V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 31. The discharge curve showing the relation between thepotential in the discharge and capacity density per gram of sulfur isrepresented by the solid line, whereas the charge curve showing therelation between the potential in the charge and capacity density pergram of sulfur is represented by the broken line.

As a result, in the test cell of comparative example 4, the initialdischarge capacity density per gram of sulfur was 1065 mAh/g which wasextremely higher than that of LiCoO₂ which has been generally used inthe positive electrode, but the discharge potential thereof was so lowas about 1.2V (vs.Li/Li⁺).

Example 12 and Example 13 show that wheretrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂and tetrahydrofuran were mixed together, the viscosity of thenon-aqueous electrolyte solution was less, which was more preferablecompared with the cases in which onlytrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂was used. Further, the results of Example 1, Example 12, Example 13, andcomparative example 4 show that where the positive electrode comprisedsulfur, the use of the mixture oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂and tetrahydrofuran is more preferable than the sole use oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂or tetrahydrofuran, according to the comparison of each dischargecapacity density around more than 2 V (vs.Li/Li⁺) in the discharge. Theamount of tetrahydrofuran therein is generally set in the range of 0.1to 99.9% by volume, preferably in the range of 0.1 to 50% by volume, andmore preferably in the range of 0.1 to 25% by volume.

EXAMPLE 14

Example 14 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in the mixturecontaining 50% by volume of 1,2-dimethoxyethane and 50% by volume oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 14.

In addition, the cyclic volutammetry of the test cell of Example 14 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.8 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 32.

As a result, in the test cell of Example 14, in the scanning in thereduction direction, the reduction peak appeared around 2.0 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation current flew around more than 2.2 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized more thansaid potential.

The test cell of Example 14 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1.0 V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 3.0V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 33. The discharge curve showing the relation between thepotential in the discharge and capacity density per gram of sulfur isrepresented by the solid line, whereas the charge curve showing therelation between the potential in the charge and capacity density pergram of sulfur is represented by the broken line.

As a result, in the test cell of Example 14, the initial dischargecapacity density per gram of sulfur was 1919 mAh/g, which is extremelyhigher than that of LiCoO₂ which has been generally used in the positiveelectrode. In addition, where both 1,2-dimethoxyethane andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂were mixed, the discharge capacity density was higher around more than1.5 V (vs.Li/Li⁺) compared with the case in which the non-aqueouselectrolyte solution comprised only 1,2-dimethoxyethane as the solventthereof as shown in the following comparative example 5, and thedischarge capacity density was also higher compared with the case inwhich the non-aqueous electrolyte solution comprised onlytrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂as the solvent thereof as shown in Example 1.

EXAMPLE 15

Example 15 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in the mixturecontaining 25% by volume of 1,2-dimethoxyethane and 75% by volume oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁻(C₃H₇)N⁻(SO₂CF₃)₂.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 15.

In addition, the cyclic volutammetry of the test cell of Example 15 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.3 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.4 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 34.

As a result, in the test cell of Example 15, in the scanning in thereduction direction, the reduction current flew around less than 2.4 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 2.5 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential.

The test cell of Example 15 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1.0 V(vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to thecharge cut-off potential of 3.0V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 35. The discharge curve showing the relation between thepotential in the discharge and capacity density per gram of sulfur isrepresented by the solid line, whereas the charge curve showing therelation between the potential in the charge and capacity density pergram of sulfur is represented by the broken line.

As a result, in the test cell of Example 15, the initial dischargecapacity density per gram of sulfur was 1636 mAh/g, which was extremelyhigher than that of LiCoO₂ which has been generally used in the positiveelectrode. In addition, where both 1,2-dimethoxyethane andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂were mixed, the discharge capacity density was higher around more than1.5 V (vs.Li/Li⁺) compared with the case in which the non-aqueouselectrolyte solution comprised only 1,2-dimethoxyethane as the solventthereof as shown in the following comparative example 5, and thedischarge capacity density was also higher compared with the case inwhich the non-aqueous electrolyte solution comprised onlytrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂as the solvent thereof as shown in Example 1.

COMPARATIVE EXAMPLE 5

Comparative example 5 used the non-aqueous electrolyte solution preparedby dissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in1,2-dimethoxyethane. Except for the above, the same procedure as that inExample 1 was taken to fabricate a test cell of comparative example 5.

In addition, the cyclic volutammetry of the test cell of comparativeexample 5 was determined as follows. The potential scanning range of thepositive electrode 11 versus the reference electrode 13 was set to be inthe range of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning ratewas set to be 1.0 mV/s. Then the operation comprising steps of scanningthe initial potential of the positive electrode 11, 2.4 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 36.

As a result, in the test cell of comparative example 5, in the scanningin the reduction direction, the reduction peak appeared around 1.8 V(vs.Li/Li⁺) and the high reduction current flew around less than 1.2 V(vs.Li/Li⁺), thus sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 2.5 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential.

The test cell of comparative example 5 was subject to the discharge atdischarging current of 0.13 mA/cm² to the discharge cut-off potential of1.0 V (vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² tothe charge cut-off potential of 3.0 V (vs.Li/Li⁺), so as to inspect theinitial charge/discharge characteristics thereof. The results are shownin FIG. 37. The discharge curve showing the relation between thepotential in the discharge and capacity density per gram of sulfur isrepresented by the solid line, whereas the charge curve showing therelation between the potential in the charge and capacity density pergram of sulfur is represented by the broken line.

As a result, in the test cell of comparative example 5, the initialdischarge capacity density per gram of sulfur was 1921 mAh/g which wasextremely higher than that of LiCoO₂ which has been generally used inthe positive electrode, but in the discharge characteristics, thecapacity density around more than 1.5 V (vs.Li/Li⁺) was less, and almostall of the discharge potential was so low as about 1.2 V (vs.Li/Li⁺).

Example 14 and Example 15 show that wheretrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂and 1,2-dimethoxyethane were mixed together, the viscosity of thenon-aqueous electrolyte solution was less, which was more preferablecompared with the cases in which onlytrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂was used. Further, the results of Example 1, Example 14, Example 15, andcomparative example 5 show that where the positive electrode comprisedsulfur, the use of the mixture oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂and 1,2-dimethoxyethane is more preferable than the sole use oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂or 1,2-dimethoxyethane, according to the comparison of each dischargecapacity density around more than 1.5 V (vs.Li/Li⁺) in the discharge.The amount of 1,2-dimethoxyethane therein is generally set in the rangeof 0.1 to 99.9% by volume, preferably in the range of 0.1 to 50% byvolume, and more preferably in the range of 0.1 to 25% by volume.

EXAMPLE 16

Example 16 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in the mixturecontaining 10% by volume of 1,3-dioxolane, 10% by volume oftetrahydrofuran, and 80% by volume oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 16.

In addition, the cyclic volutammetry of the test cell of Example 16 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.8 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 38.

As a result, in the test cell of Example 16, in the scanning in thereduction direction, the reduction peak appeared around 1.9 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 2.4 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential. The same results are attained also from the second cycle andonwards, and it was proved that sulfur reversibly reacted around 2 V(vs.Li/Li⁺).

In the mixture of 1,3-dioxolane, tetrahydrofuran, andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂,the ratio of trimethylpropylammonium.bis(trifluoromethylsulfonyl)imidetherein is generally set in the range of 0.1 to 99.9% by volume,preferably in the range of 50 to 99.9% by volume, and more preferably inthe range of 80 to 99.9% by volume.

EXAMPLE 17

Example 17 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in the mixturecontaining 10% by volume of 1,3-dioxolane, 10% by volume of1,2-dimethoxyethane, and 80% by volume oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂.Except for the above, the same procedure as that in Example 1 was takento fabricate a test cell of Example 17.

In addition, the cyclic volutammetry of the test cell of Example 17 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.7 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 39.

As a result, in the test cell of Example 17, in the scanning in thereduction direction, the reduction peak appeared around 1.9 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 2.5 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential. The same results were attained also from the second cycle andonwards, and it was proved that sulfur reversibly reacted around 2 V(vs.Li/Li⁺).

In the mixture of 1,3-dioxolane, 1,2-dimethoxyethane, andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂,the ratio of trimethylpropylammonium.bis(trifluoromethylsulfonyl)imidetherein is generally set in the range of 0.1 to 99.9% by volume,preferably in the range of 50 to 99.9% by volume, and more preferably inthe range of 80 to 99.9% by volume.

EXAMPLE 18

Example 18 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in the mixturecontaining 1,3-dioxolane, tetrahydrofuran, 1,2-dimethoxyethane, andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂in a volume ratio of 6.7:6.7:6.7:80. Except for the above, the sameprocedure as that in Example 1 was taken to fabricate a test cell ofExample 18.

In addition, the cyclic volutammetry of the test cell of Example 18 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.85 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 40.

As a result, in the test cell of Example 18, in the scanning in thereduction direction, the reduction peak appeared around 1.9 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 2.5 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential. The same results were attained also from the second cycle andonwards, and it is believed that sulfur reversibly reacted around 2 V(vs.Li/Li⁺).

In the mixture of 1,3-dioxolane, tetrahydrofuran, 1,2-dimethoxyethane,andtrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide(CH₃)₃N⁺(C₃H₇)N⁻(SO₂CF₃)₂,the ratio of trimethylpropylammonium.bis(trifluoromethylsulfonyl)imidetherein is generally set in the range of 0.1 to 99.9% by volume,preferably in the range of 50 to 99.9% by volume, and more preferably inthe range of 80 to 99.9% by volume.

EXAMPLE 19

Example 19 used the non-aqueous electrolyte solution prepared bydissolving 0.5 mol/l of LiN(CF₃SO₂)₂ as the lithium salt in the mixturecontaining 50% by volume of 1,3-dioxolane and 50% by volume of1-ethyl-3-methylimidazolium bis(pentafluoroethylsulfonyl)imide(C₂H₅)(C₃H₃N₂)⁺(CH₃)N⁻(SO₂C₂F₅)₂. Exceptfor the above, the same procedure as that in Example 1 was taken tofabricate a test cell of Example 19.

In addition, the cyclic volutammetry of the test cell of Example 19 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 3.0 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.3 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 41.

As a result, in the test cell of Example 19, in the scanning in thereduction direction, the reduction peak appeared around 1.8 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 2.7 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential.

The test cell of Example 19 was subject to the discharge at dischargingcurrent of 0.13 mA/cm² to the discharge cut-off potential of 1.0 V (vs.Li/Li⁺) and the charge at charging current of 0.13 mA/cm² to the chargecut-off potential of 3.0V (vs.Li/Li⁺), so as to inspect the initialcharge/discharge characteristics thereof. The results are shown in FIG.42. The discharge curve showing the relation between the potential inthe discharge and capacity density per gram of sulfur is represented bythe solid line, whereas the charge curve showing the relation betweenthe potential in the charge and capacity density per gram of sulfur isrepresented by the broken line.

As a result, in the test cell of Example 19, the initial dischargecapacity density per gram of sulfur was 741 mAh/g which was extremelyhigher than that of LiCoO₂ which has been generally used in the positiveelectrode. Further, where 1,3-dioxolane and 1-ethyl-3-methylimidazoliumbis (pentafluoroethylsulfonyl) imide(C₂H₅)(C₃H₃N₂)⁺(CH₃)N⁻(SO₂C₂F₅)₂were mixed, the capacity density in the discharge around more than 2.0 V(vs.Li/Li⁺) was higher compared with the case in which only1,3-dioxolane was used.

However, the discharge capacity density thereof was less than those ofExamples 10 and 11.

EXAMPLE 20

Example 20 used the non-aqueous electrolyte solution prepared bydissolving 1 mol/l of tetramethylammonium.tetrafluoroborate (CH₃)₄N⁺BF₄⁻ as quaternary ammonium salt and 1 mol/l of LiPF₆ as the lithium saltin a mixed solvent containing ethylene carbonate and diethyl carbonatein a volume ratio of 1:1. Except for the above, the same procedure asthat in Example 1 was taken to fabricate a test cell of Example 20.

In addition, the cyclic volutammetry of the test cell of Example 20 wasdetermined as follows. The potential scanning range of the positiveelectrode 11 versus the reference electrode 13 was set to be in therange of 1.0 to 4.7 V (vs.Li/Li⁺), while the potential scanning rate wasset to be 1.0 mV/s. Then the operation comprising steps of scanning theinitial potential of the positive electrode 11, 2.7 V (vs.Li/Li⁺),versus the reference electrode 13 in the reduction direction, andscanning said potential in the oxidation direction was carried out for 3cycles, so as to determine the cyclic volutammetry at each cycle. Theresults are shown in FIG. 43.

As a result, in the test cell of Example 20, in the scanning in thereduction direction, the reduction peak appeared around 1.6 V(vs.Li/Li⁺), and sulfur was expectedly reduced. In the scanning in theoxidation direction, the oxidation peak appeared around 4.4 V(vs.Li/Li⁺), and the reduced sulfur was expectedly oxidized around saidpotential. In the second and third cycle, the oxidation and thereduction peaks were less than those in the first cycle, but theoxidation and the reduction peaks appeared in the same way, therefore,it is believed that sulfur reversibly reacted.

INDUSRICAL APPLICABILITY

As described above, in the invention, a non-aqueous electrolytesecondary cell is provided with a positive electrode comprising only asimple substance of sulfur as an active material, wherein a non-aqueouselectrolyte solution comprises any of a room-temperature molten salthaving a melting point of 60° C. or less, the room-temperature moltensalt having the melting point of 60° C. or less and lithium salt, theroom-temperature molten salt having the melting point of 60° C. or lessand at least one solvent selected from circular ether, chain ether, andcarbonate fluoride, or a quaternary ammonium salt and a lithium salt, sothat sulfur in the positive electrode reversibly reacts with lithiumeven in normal temperature and the non-aqueous electrolyte secondarycell is capable of charge/discharge reactions in the normal temperature.

In addition, where the non-aqueous electrolyte solution comprises theroom-temperature molten salt having the melting point of 60° C. or lessand sulfur reduction product, sulfur reversibly reacts with lithium inthe positive electrode in the normal temperature, the non-aqueouselectrolyte secondary cell is capable of charge/discharge reactions inthe normal temperature, and even in a case in which the positiveelectrode comprises sulfur, the non-aqueous electrolyte secondary cellis charged/discharged in the normal temperature.

In the non-aqueous electrolyte secondary cell according to the presentinvention, where the positive electrode comprises sulfur, capacity perunit weight is further improved compared with the case in which anorganic disulfide compound is used.

1. A non-aqueous electrolyte secondary cell provided with a positiveelectrode, a negative electrode, and a non-aqueous electrolyte solution,wherein said positive electrode comprises elemental sulfur as an activematerial and said non-aqueous electrolyte solution comprises aroom-temperature molten salt having a melting point of 60° C. or less,wherein said room-temperature molten salt having the melting point of60° C. or less in the non-aqueous electrolyte solution is a quaternaryammonium salt.
 2. The non-aqueous electrolyte secondary cell as claimedin claim 1, wherein said negative electrode comprises material capableof absorbing and desorbing lithium.
 3. The non-aqueous electrolytesecondary cell as claimed in claim 1, wherein said negative electrodecomprises carbon material or silicon material.
 4. The non-aqueouselectrolyte secondary cell as claimed in claim 1, wherein saidquaternary ammonium salt is at least one oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide,trimethyloctylammonium.bis(trifluoromethylsulfonyl)imide,trimethylallylammonium.bis(trifluoromethylsulfonyl)imide,trimethylhexylammonium.bis(trifluoromethylsulfonyl)imide,trimethylethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,trimethylallylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,trimethylpropylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,tetraethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,andtriethylmethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide.5. The non-aqueous electrolyte secondary cell as claimed in claim 1,wherein said non-aqueous electrolyte solution comprises 50% by volume ormore of the room-temperature molten salt having the melting point of 60°C. or less as a solvent thereof.
 6. The non-aqueous electrolytesecondary cell as claimed in claim 1, wherein said non-aqueouselectrolyte solution comprises the room-temperature molten salt havingthe melting point of 60° C. or less and a lithium salt.
 7. Thenon-aqueous electrolyte secondary cell as claimed in claim 1, whereinconductive agent is added to said positive electrode.
 8. A non-aqueouselectrolyte secondary cell provided with a positive electrode, anegative electrode, and a non-aqueous electrolyte solution, wherein saidnon-aqueous electrolyte solution comprises a room-temperature moltensalt having a melting point of 60° C. or less and sulfur reductionproduct; and wherein said positive electrode comprises elemental sulfuras an active material.
 9. The non-aqueous electrolyte secondary cell asclaimed in claim 8, wherein said negative electrode comprises materialcapable of absorbing and desorbing lithium.
 10. The non-aqueouselectrolyte secondary cell as claimed in claim 8, wherein said negativeelectrode comprises carbon material or silicon material.
 11. Thenon-aqueous electrolyte secondary cell as claimed in claim 8, whereinsaid room-temperature molten salt having the melting point of 60° C. orless in the non-aqueous electrolyte solution is a quaternary ammoniumsalt.
 12. The non-aqueous electrolyte secondary cell as claimed in claim11, wherein said quaternary ammonium salt is at least one oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide,trimethyloctylammonium.bis(trifluoromethylsulfonyl)imide,trimethylallylammonium.bis(trifluoromethylsulfonyl)imide,trimethylhexylammonium.bis(trifluoromethylsulfonyl)imide,trimethylethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)trimethylallylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,trimethylpropylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,tetraethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,andtriethylmethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide.13. The non-aqueous electrolyte secondary cell as claimed in claim 9,wherein said sulfur reduction product is elemental sulfur reduced in theroom-temperature molten salt having the melting point of 60° C. or less.14. A non-aqueous electrolyte secondary cell provided with a positiveelectrode, a negative electrode, and a non-aqueous electrolyte solution,wherein said positive electrode comprises elemental sulfur as an activematerial and said non-aqueous electrolyte solution comprises aroom-temperature molten salt having a melting point of 60° C. or lessand at least one solvent selected from cyclic ether, chain ether, andcarbonate fluoride; wherein said room-temperature molten salt having themelting point of 60° C. or less in the non-aqueous electrolyte solutionis a quaternary ammonium salt.
 15. The non-aqueous electrolyte secondarycell as claimed in claim 14, wherein said negative electrode comprisesmaterial capable of absorbing and desorbing lithium.
 16. The non-aqueouselectrolyte secondary cell as claimed in claim 14, wherein said negativeelectrode comprises carbon material or silicon material.
 17. Thenon-aqueous electrolyte secondary cell as claimed in claim 16, whereinsaid non-aqueous electrolyte solution comprises 50% by volume or more ofthe room-temperature molten salt having the melting point of 60° C. orless as a solvent thereof.
 18. The non-aqueous electrolyte secondarycell as claimed in claim 14, wherein said quaternary ammonium salt is atleast one of trimethylpropylammonium.bis(trifluoromethylsulfonyl)imide,trimethyloctylammonium.bis(trifluoromethylsulfonyl)imide,trimethylallylammonium.bis(trifluoromethylsulfonyl)imide,trimethylhexylammonium.bis(trifluoromethylsulfonyl)imide,trimethylethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,trimethylallylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,trimethylpropylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,tetraethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,andtriethylmethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide.19. The non-aqueous electrolyte secondary cell as claimed in claim 14,wherein said cyclic ether is at least one of 1,3-dioxolane,4-methyl-1,3-dioxolane, tetrahydrofuran, 2-methyltetrahydrofuran,propylene oxide, 1,2-butylene oxide, 1,4-dioxane, 1,3,5-trioxane, furan,2-methylfuran, 1,8-cineol, and crown ether, said chain ether is at leastone of 1,2-dimethoxyethane, diethyl ether, dipropyl ether, diisopropylether, dibuthyf ether, dihexyl ether, ethylvinyl ether, buthylvinylether, methyiphenyl ether, ethyiphenyl ether, buthylphenyl ether,pentyiphenyl ether, methoxy toluene, benzylethyl ether, diphenyl ether,dibenzyl ether, o-dimethoxy benzene, 1,2-diethoxyethane,1,2-dibutoxyethane, diethylene glycol dimethyl ether, diethylene glycoldiethyl ether, diethylene glycol dibuthyl ether, 1,1-dimethoxy methane,1,1-diethoxy ethane, triethylene glycol dimethyl ether, andtetraethylene glycol dimethyl ether, and said carbonate fluoride is atleast one of trifluoropropylene carbonate and fluoroethylene carbonate.20. A non-aqueous electrolyte secondary cell provided with a positiveelectrode, a negative electrode comprising material capable of absorbingand desorbing lithium, and a non-aqueous electrolyte solution, whereinsaid positive electrode comprises only elemental sulfur as an activematerial and said non-aqueous electrolyte solution comprises aquaternary ammonium salt and a lithium salt.
 21. The non-aqueouselectrolyte secondary cell as claimed in claim 20, wherein saidquaternary ammonium salt is at least one oftrimethylpropylammonium.bis(trifluoromethylsulfonyl)imide,trimethyloctylammonium.bis(trifluoromethylsulfonyl)imide,trimethylallylammonium.bis(trifluoromethylsulfonyl)imide,trimethylhexylammonium.bis(trifluoromethylsulfonyl)imide,trimethylethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,trimethylallylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,trimethylpropylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,tetraethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide,triethylmethylammonium.2,2,2-trifluoro-N-(trifluoromethylsulfonyl)acetamide, tetramethylammonium.tetrafluoroborate,tetramethylammonium.hexafluorophosphate,tetraethylammonium.tetrafluoroborate, andtetraethylammonium.hexafluorophosphate.